Chemistry: Matter and ChangeSection 10.1 Measuring Matter Section 10.2 Mass and the Mole Section...
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Chapter 10: The Mole CHEMISTRY Matter and Change
Transcript of Chemistry: Matter and ChangeSection 10.1 Measuring Matter Section 10.2 Mass and the Mole Section...
Chapter 10: The Mole
CHEMISTRY Matter and Change
Section 10.1 Measuring Matter
Section 10.2 Mass and the Mole
Section 10.3 Moles of Compounds
Section 10.4 Empirical and Molecular Formulas
Section 10.5 Formulas of Hydrates
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Table Of Contents CHAPTER
10
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• Explain how a mole is used to indirectly count the number of particles of matter.
molecule: two or more atoms that covalently bond together to form a unit
mole
Avogadro’s number
• Relate the mole to a common everyday counting unit.
• Convert between moles and number of representative particles.
Chemists use the mole to count atoms, molecules, ions, and formula units.
SECTION
10.1 Measuring Matter
Counting Particles • Chemists need a convenient method for accurately
counting the number of atoms, molecules, or formula units of a substance.
• The mole is the SI base unit used to measure the amount of a substance.
• 1 mole is the amount of atoms in 12 g of pure carbon-12, or 6.02 1023 representative particles, which is any kind of particle – an atom, a molecule, a formula unit, an electron, an ion, etc.
• The number is called Avogadro’s number.
SECTION
10.1 Measuring Matter
Converting Between Moles and Particles
• Conversion factors must be used.
• Moles to particles
Number of molecules in 3.50 mol of sucrose
SECTION
10.1 Measuring Matter
Converting Between Moles and Particles (cont.)
• Particles to moles
• Use the inverse of Avogadro’s number as the conversion factor.
SECTION
10.1 Measuring Matter
What does the mole measure?
A. mass of a substance
B. amount of a substance
C. volume of a gas
D. density of a gas
SECTION
10.1 Section Check
D. 1 mol 6.02 1023 particles
SECTION
10.1 Section Check
What is the conversion factor for
determining the number of moles of a
substance from a known number of
particles?
A.
C. 1 particle 6.02 1023
B.
• Relate the mass of an atom to the mass of a mole of atoms.
conversion factor: a ratio of equivalent values used to express the same quantity in different units
molar mass
• Convert between number of moles and the mass of an element.
• Convert between number of moles and number of atoms of an element.
A mole always contains the same number of particles; however, moles of different substances have different masses.
SECTION
10.2 Mass and the Mole
The Mass of a Mole
• 1 mol of copper (6.02 x 1023 atoms of copper) and 1 mol of carbon (6.02 x 1023 atoms of carbon) have different masses.
• One copper atom has a different mass than 1 carbon atom.
SECTION
10.2 Mass and the Mole
The Mass of a Mole (cont.)
• Molar mass is the mass in grams of one mole of any pure substance.
• The molar mass of any element is numerically equivalent to its atomic mass and has the units g/mol.
SECTION
10.2 Mass and the Mole
Using Molar Mass
• Moles to mass
3.00 moles of copper has a mass of 191 g.
SECTION
10.2 Mass and the Mole
Using Molar Mass (cont.)
• Convert mass to moles with the inverse molar mass conversion factor.
• Convert moles to atoms with Avogadro’s number as the conversion factor.
SECTION
10.2 Mass and the Mole
Using Molar Mass (cont.)
• This figure shows the steps to complete conversions between mass and atoms.
SECTION
10.2 Mass and the Mole
The mass in grams of 1 mol of any pure substance is:
A. molar mass
B. Avogadro’s number
C. atomic mass
D. 1 g/mol
SECTION
10.2 Section Check
Molar mass, in its un-inverted form, is used to convert what?
A. mass to moles
B. moles to mass
C. atomic weight
D. particles
SECTION
10.2 Section Check
• Recognize the mole relationships shown by a chemical formula.
representative particle: an atom, molecule, formula unit, or ion
• Calculate the molar mass of a compound.
• Convert between the number of moles and mass of a compound.
• Apply conversion factors to determine the number of atoms or ions in a known mass of a compound.
SECTION
10.3 Moles of Compounds
The molar mass of a compound can be calculated from its chemical formula and can be used to convert from mass to moles of that compound.
SECTION
10.3 Moles of Compounds
Chemical Formulas and the Mole
• Chemical formulas indicate the numbers and types of atoms contained in one unit of the compound.
• One mole of CCl2F2 contains one mole of C atoms, two moles of Cl atoms, and two moles of F atoms.
SECTION
10.3 Moles of Compounds
The Molar Mass of Compounds
• The molar mass of a compound equals the molar mass of each element, multiplied by the moles of that element in the chemical formula, added together.
• The molar mass of a compound demonstrates the law of conservation of mass.
SECTION
10.3 Moles of Compounds
Converting Moles of a Compound to Mass
• For elements, the conversion factor is the molar mass of the compound.
• The procedure is the same for compounds, except that you must first calculate the molar mass of the compound.
SECTION
10.3 Moles of Compounds
Converting the Mass of a Compound to
Moles
• The conversion factor is the inverse of the molar mass of the compound.
SECTION
10.3 Moles of Compounds
Converting the Mass of a Compound to
Number of Particles
• Convert mass to moles of compound with the inverse of molar mass.
• Convert moles to particles with Avogadro’s number.
SECTION
10.3 Moles of Compounds
Converting the Mass of a Compound to
Number of Particles (cont.)
• This figure summarizes the conversions between mass, moles, and particles.
SECTION
10.3 Moles of Compounds
How many moles of OH— ions are in 2.50 moles of Ca(OH)2?
A. 2.00
B. 2.50
C. 4.00
D. 5.00
SECTION
10.3 Section Check
How many particles of Mg are in 10 moles of MgBr2?
A. 6.02 1023
B. 6.02 1024
C. 1.20 1024
D. 1.20 1025
SECTION
10.3 Section Check
• Explain what is meant by the percent composition of a compound.
percent by mass: the ratio of the mass of each element to the total mass of the compound expressed as a percent
percent composition
empirical formula
molecular formula
• Determine the empirical and molecular formulas for a compound from mass percent and actual mass data.
A molecular formula of a compound is a whole-number multiple of its empirical formula.
SECTION
10.4 Empirical and Molecular Formulas
Percent Composition
• The percent by mass of any element in a compound can be found by dividing the mass of the element by the mass of the compound and multiplying by 100.
SECTION
10.4 Empirical and Molecular Formulas
Percent Composition (cont.)
• The percent by mass of each element in a compound is the percent composition of a compound.
• Percent composition of a compound can also be determined from its chemical formula.
SECTION
10.4 Empirical and Molecular Formulas
Empirical Formula
• The empirical formula for a compound is the smallest whole-number mole ratio of the elements.
• You can calculate the empirical formula from percent by mass by assuming you have 100.00 g of the compound. Then, convert the mass of each element to moles.
• The empirical formula may or may not be the same as the molecular formula.
Molecular formula of hydrogen peroxide = H2O2
Empirical formula of hydrogen peroxide = HO
SECTION
10.4 Empirical and Molecular Formulas
Molecular Formula
• The molecular formula specifies the actual number of atoms of each element in one molecule or formula unit of the substance.
• Molecular formula is always a whole-number multiple of the empirical formula.
SECTION
10.4 Empirical and Molecular Formulas
Molecular Formula (cont.)
SECTION
10.4 Empirical and Molecular Formulas
What is the empirical formula for the compound C6H12O6?
A. CHO
B. C2H3O2
C. CH2O
D. CH3O
SECTION
10.4 Section Check
Which is the empirical formula for hydrogen peroxide?
A. H2O2
B. H2O
C. HO
D. none of the above
SECTION
10.4 Section Check
• Explain what a hydrate is and relate the name of the hydrate to its composition.
crystal lattice: a three-dimensional geometric arrangement of particles
hydrate
• Determine the formula of a hydrate from laboratory data.
Hydrates are solid ionic compounds in which water molecules are trapped.
SECTION
10.5 Formulas of Hydrates
Naming Hydrates
• A hydrate is a compound that has a specific number of water molecules bound to its atoms.
• The number of water molecules associated with each formula unit of the compound is written following a dot.
• Sodium carbonate decahydrate = Na2CO3 • 10H2O
SECTION
10.5 Formulas of Hydrates
Analyzing a Hydrate
• When heated, water molecules are released from a hydrate leaving an anhydrous compound.
• To determine the formula of a hydrate, find the number of moles of water associated with 1 mole of hydrate.
SECTION
10.5 Formulas of Hydrates
Analyzing a Hydrate (cont.)
• Weigh hydrate.
• Heat to drive off the water.
• Weigh the anhydrous compound.
• Subtract and convert the difference to moles.
• The ratio of moles of water to moles of anhydrous compound is the coefficient for water in the hydrate.
SECTION
10.5 Formulas of Hydrates
Use of Hydrates
• Anhydrous forms of hydrates are often used to absorb water, particularly during shipment of electronic and optical equipment.
• In chemistry labs, anhydrous forms of hydrates are used to remove moisture from the air and keep other substances dry.
SECTION
10.5 Formulas of Hydrates
Heating a hydrate causes what to happen?
A. Water is driven from the hydrate.
B. The hydrate melts.
C. The hydrate conducts electricity.
D. There is no change in the hydrate.
SECTION
10.5 Section Check
A hydrate that has been heated and the water driven off is called:
A. dehydrated compound
B. antihydrated compound
C. anhydrous compound
D. hydrous compound
SECTION
10.5 Section Check
Chemistry Online
Study Guide
Chapter Assessment
Standardized Test Practice
Study Guide
CHAPTER
10 The Mole
Key Concepts
• The mole is a unit used to count particles of matter indirectly. One mole of a pure substance contains Avogadro’s number of particles.
• Representative particles include atoms, ions, molecules, formula units, electrons, and other similar particles.
• One mole of carbon-12 atoms has a mass of exactly 12 g.
• Conversion factors written from Avogadro’s relationship can be used to convert between moles and number of representative particles.
SECTION
10.1 Measuring Matter
Study Guide
Key Concepts
• The mass in grams of 1 mol of any pure substance is called its molar mass.
• The molar mass of an element is numerically equal to its atomic mass.
• The molar mass of any substance is the mass in grams of Avogadro’s number of representative particles of the substance.
• Molar mass is used to convert from moles to mass. The inverse of molar mass is used to convert from mass to moles.
SECTION
10.2 Mass and the Mole
Study Guide
Key Concepts
• Subscripts in a chemical formula indicate how many moles of each element are present in 1 mol of the compound.
• The molar mass of a compound is calculated from the molar masses of all of the elements in the compound.
• Conversion factors based on a compound’s molar mass are used to convert between moles and mass of a compound.
SECTION
10.3 Moles of Compounds
Study Guide
Key Concepts
• The percent by mass of an element in a compound gives the percentage of the compound’s total mass due to that element.
• The subscripts in an empirical formula give the smallest whole-number ratio of moles of elements in the compound.
• The molecular formula gives the actual number of atoms of each element in a molecule or formula unit of a substance.
• The molecular formula is a whole-number multiple of the empirical formula.
SECTION
10.4 Empirical and Molecular Formulas
Study Guide
Key Concepts
• The formula of a hydrate consists of the formula of the ionic compound and the number of water molecules associated with one formula unit.
• The name of a hydrate consists of the compound name and the word hydrate with a prefix indicating the number of water molecules in 1 mol of the compound.
• Anhydrous compounds are formed when hydrates are heated.
SECTION
10.5 Formulas of Hydrates
Study Guide
What does Avogadro’s number represent?
A. the number of atoms in 1 mol of an element
B. the number of molecules in 1 mol of a compound
C. the number of Na+ ions in 1 mol of NaCl (aq)
D. all of the above
The Mole
Chapter Assessment
CHAPTER
10
The molar mass of an element is numerically
equivalent to what?
A. 1 amu
B. 1 mole
C. its atomic mass
D. its atomic number
Chapter Assessment
The Mole CHAPTER
10
How many moles of hydrogen atoms are in one mole of H2O2?
A. 1
B. 2
C. 3
D. 0.5
Chapter Assessment
The Mole CHAPTER
10
What is the empirical formula of Al2Br3?
A. AlBr
B. AlBr3
C. Al2Br
D. Al2Br3
Chapter Assessment
The Mole CHAPTER
10
What is an ionic solid with trapped water molecules called?
A. aqueous solution
B. anhydrous compound
C. hydrate
D. solute
Chapter Assessment
The Mole CHAPTER
10
Two substances have the same percent by mass composition, but very different properties. They must have the same ____.
A. density
B. empirical formula
C. molecular formula
D. molar mass
Standardized Test Practice
The Mole CHAPTER
10
How many moles of Al are in 2.0 mol of Al2Br3?
A. 2
B. 4
C. 6
D. 1
Standardized Test Practice
The Mole CHAPTER
10
How many water molecules are associated with 3.0 mol of CoCl2 • 6H2O?
A. 18
B. 1.1 1025
C. 3.6 1024
D. 1.8 1024
Standardized Test Practice
The Mole CHAPTER
10
How many atoms of hydrogen are in 3.5 mol of H2S?
A. 7.0 1023
B. 2.1 1023
C. 6.0 1023
D. 4.2 1024
The Mole CHAPTER
10
Standardized Test Practice
Which is not the correct formula for an ionic compound?
A. CO2
B. NaCl
C. Na2SO4
D. LiBr2
The Mole CHAPTER
10
Standardized Test Practice
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